Bonding to Hydrogen

A Lousy Acid, a Lousy Base

Molecular hydrogen is pretty unreactive, as is methane. Hydrogen burns, of course (with a flame that is nearly colorless but very, very hot). But to get it to burn you need a match, even though the reaction to form water, Cavendish’s and Lavoisier’s reaction, gives off ~68 kcal/mol of dihydrogen burned. That’s chemistry: Things that should spontaneously proceed by the dictates of thermodynamics (like hydrogen burning) actually encountering substantial barriers to doing so.

Chemical reactivity is predominantly that of acids and bases—that is why we spend so much time in introductory chemistry on this property of molecules. A base (ammonia, for example) is a good donor of electrons; in MO terms it has an energetically high-lying filled molecular orbital. An acid (the hydronium ion, the aquated proton, H3O+) is a good acceptor of electrons, as it has a low-energy empty MO. Hydrogen has an occupied MO, just one; you’ve seen it—it’s the σg in the MO picture of the molecule (see figure at left top). That MO lies low in energy; H2’s ionization potential, a measure of the energy of that MO, is large, 15.4 eV. And H2’s lowest unoccupied MO, σu*, is relatively high lying—to promote an electron from the filled MO to the unfilled one takes ~11 eV.

Put into plain English, the hydrogen molecule is a lousy base and a lousy acid. The molecule is then relatively unreactive, even as it burns giving off a good bit of heat. Other molecules lack a good handhold, so to speak, on H2.